Based on the Gibbs free energy equation (ΔG = ΔH − TΔS), which of the following correctly explains when a reaction might be endothermic but exergonic?

An endothermic reaction becomes exergonic only when the endothermic reaction involves an entropy increase large enough to cause the (−TΔS) term in the ΔG equation to outweigh the positive ΔH term, resulting in a negative ΔG.

An endothermic reaction becomes exergonic only when the endothermic process involves an entropy decrease large enough to cause the (−TΔS) term in the ΔG equation to add to the positive ΔH term, resulting in a positive ΔG.

An endothermic reaction becomes exergonic when the endothermic reaction involves an entropy increase small enough to cause the (−TΔS) term in the ΔG equation to outweigh the positive ΔH term, resulting in a negative ΔG.

An endothermic reaction cannot be an exergonic reaction because ΔH is positive, which by default, results in positive ΔG.